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DETERMINATION OF THE ELECTRODE POTENTIAL OF CHROMIUM 


MELVILLE YORK ENYART 


THESIS 
FOR THE 


DEGREE OF BACHELOR OF SCIENCE 


CHEMISTRY 


COLLEGE OF LIBERAL ARTS AND SCIENCES 
UNIVERSITY OF ILLINOIS 


1921 


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TABLE OF CONTENTS 


PaCMe a Wie GRO. ea ad én. BHAS. aid de. Page 1 
TRCOGES EL CALS, 0.6 wine ashe 

Historical. 

Experimental 


SWIM LY, « ace, sw oan0 


Bibliography 


Acknowledgment 


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(I) 


INTRODUCTION 


The importance of an eccurate knowledge of the potential devel- 
oped between a metal and an aqueous solution of a salt of that metal 
becomes apparent at once when a study of the common cells is begun. 
From such study we come to realize that aside from the purely scien- 
tific or theoretical standpoint of the question there is a practical 
side as well. 

In orderto assure the proper placing of the elements in the 
electromotive series an accurate determination of the electrode poter 
tials is necessary. Many of these have already been accurately deter 
mined,some by direct and others by indirect methods; the latter 
based upon assumptions that may or may not be valid. 

.Among those metals that have not been accurately determined is 
chromium. It was therefore suggested to the author that such a deter 
mination would prove of value,as mentioned above, both from a pure 
science and practical viewpoint. With these things in mind the 
author set out to make such a determination,if it were possible. 

In order that sufficient experience to handle the admittedly 
more complex problem of determining the potential of chromium be 


obtained, it was decided to check up certain values as obtained 


' 
by recent workers upon zinc. 


The results thus obtained were not just what it was hoped they 
might be although the apparatus as used by Mr. David Schlesinger 
and the author did not have some of the refinements that the other 
workers had. That is,to say, tnis apparatus lacked for one thing a 


thermostatic control of temperature. This was in ace rtain measure 


unfair to the previous work that had been done,but it was desired 


Digitized by the Internet Archive 
in 2016 


https://archive.org/details/determinationofeOOenya 


(2) 
to obtain some data that might be used in the solution of a problem 
connected with a laboratory course in which the refinements of ther- 


mostatic control could not be used to advantage. After the suggestioy 


in regard to the use of finely divided zinc prepared electrolytically 


had been carried out, the attention of the author was again turned 
toward the real problem,that is the determination of the specific 


electrode potential of chromium. 


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(3) 
THEORETICAL 


The question of the source of the electromotive force of a cell 
arouses our interest in electro-chemical processes. A complete 
answer to this question involves a study of some of the simpler cells 
from the theoretical side. Those in whicn we shall be interested 
are those involving electro-chemical, rather than physical processes. 
The common Daniell cell is familiar to most of us and isa 
good example of what we mean by a cell. Any sucn arrangement for 


performing electro-chemical reactions is know as a cell. While the 


general form of cells is the same,there are are two classes into whick 
they may be divided,namely,electrolytic and voltaic. We find then in 
analyzing such a cell that it consists of,(a) a zine electrode in a 
solution of zine sulfate,(b)a copper electrode in a solution of 
copper sulfate from which the entire cell is formed poy placing the 

two half cells in conducing contact. Such a combination will yield 

@ current under suitable conditions,that is, by joining tne terminal s| 
of the electrodes by a conductor. 

When a change in state can be made to take place in some such 
manner as above in which reduction takes place at one electrode and 
oxidation at the other,,it can be made to yield electrical energy. 
Reduction processes are connected with the giving up of electricity 
and oxidation processes with the taking on of charges. Whenever 
these two actions can be separated but yet be in electrical contact 
with one another, we have a cell. 

The quantity of work that the system is capapdle of giving depend§ 
upon the initial and final states of the system. This then,involves 


the temperature, the concentration, the pressures, and the chemical 


reactions taking place. It is to be noted then that ‘the driving 


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(4) 

force or tendency toward chemical reaction in a cell is measured by 
the voltage or electromotive force that is developed. The work done 
per unit quantity of electricity flowing througn the cell may be 
given as a definition of this quantity. 

This then,shows how closely related are the transformations 
of energy accompanying a chemical reaction and the electromotive 
force it is capable of producing,if carried out in a cell, and that 
this force may be regarded as a measure to a certain extent of the 
energy change. 

The electromotive force of a cell is made up of the sum of the 
forces produced at the junctions of dissimilar substances. Thus the 
Plectromotive forceof the Daniell cell is the algebraic sum of the 
force from the zine to the zine sulfate, that from the zine sulfate 
solution to the copper sulfate solution, and that from the copper 
sulfate solution to the copper. The electromotive force produced at 
the junction of the two solutions is known as the liquid potential 
and that produced at the electrodes as tne electrode potential. 

While certain attempts at determination of absolute values of 
electrode potentials have been made,sucn as by the use of the 


dropping electrode method and the capillary electrometer, these have 


not been very successful. That is,to say, the values so far obtained : 


ere not as good as the comparative values that are always meant 
unless otherwise specified. Measurements made by the use of ordinary 
half cells as a basis of reference are more easily made,and from 
such measurements our data relative to the electrode potentials have 
been compiled. There are several standard reference electrodes in use 
and several scales are employed but we shall make use of the normal 


hydrogen electrode in all our considerations in which the potential 


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(5) 
developed between a platinum electrode saturated with hydrogen gas 
under a pressure of one atmosphere dipping into a solution containing 
hydrogen ion at a concentration of one formula weight per liter is 
placed equal to zero. 
From a consideration of the change in state which a cell under- 


goes there 1s developed the expression for the free energy decrease 


experienced in the process:- 


~aF=NAT A & 
2 


in which c, and c,are tne initial and final concentrations, R is the 
gas constant, and T is the absolute temperature,while the log term 
lis to the base e. With the aid of the other expression for tne free 
energy decrease:- 

Sa LV 
we are able to arrive at an expression whereby the change in tne elec 


tromotive force developed by a given change can be calculateds- 


nF & 
Making use of this relation for the changé in free energy upon dilu- 
tion, assuming that the solution is completely ionized and follows 
precisely the gas laws, Nernst arrived at the well known and importan 
formula wnich goes by his name namely:- 

2 = Le + 2.5026 x AL /0fe 

where Eis the potential difference for aus gee cohtontraiwen C; 
E,is the electrode potential in the case where C,has the value unity: 
R is the gas constant with a value of 8.32 joules per degree; n is th 
valence of the ion; ¥ is the farady or 96,000 coulombs. This gives 
then a method of determining the electrode potential #,,if we know 


the other members of the equation. 


Due to the very nature of some of tne elements it is quite easil 


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(6) 

seen that the use of a simple aqueous solution of a salt of the el- 
ement whose electrode potential is to be measured combined with an 
electrode of this material is out of the question. Sone of the metals 
exhibit such properties as passivity that quite materially affect 
an accurate determination of potentials » y such a method. There 
have been as a result of these difficulties some very ingenious 
devices and methods brought forth for getting around them. It may be 
well at this time to pay some attention to a few of these that we 
may have before us the possinilities in the way of methods that have 
been tried. 

Aluminum,as is well known,exhibitsthe tendency toward passivity. 

In attempting to get a value for the electrode potential of this 
metal Kahlenburg~ found tha the effect of a solution of sodium 
hydroxide was the same as amalgamation of the electrode. That is,to 
say,that in so far as the caustic solution acts as a solvent for the 
oxide of the metal and tnereby removes tne tendency toward passivity 
the mercury is thought to do likewise. The use of amalgams we find 
has been fairly general even in the work with non-aqueous solutions. 


The potential developed by lithium was determined by Noyes and 


3 
Keyes by the use. of the systems LisLiCl + propylamine/Li-Hg(0.035 ~~) 


and Li-Hg LiOH/LiC1/KCl/nKCl nHgCl. fn using such a dilute amalgam 
the action of the water upon the lithium was practically nil while 
in the propylamine there was no action. Then upon the assumption that 
the difference infpotential betweenthe lithium and the lithium amalgam 
in propylamine is the same as would be given by tne two in an aqueous 
solution they were able to get a value for the electrode potential 
of lithium. Similar methods have been used for other very reactive 


metals. 
It is interesting to note that the potential of zinc,copper, 


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4 
and silver have been determined by electrostatic measurements. In 


this method a glass flask was covered with a layer of the metal 
and then a knowncharge applied. By means of an electrometer the 
values were determinéd for solutions of fairly high concentration, 


that is, normal and half normal. 


Working upon the assumption,based upon theoretical consideratiogs 
,that the ditference between the electrode potential of a metal in 
ethyl alcohol and water for all metals has the same value,Roderburg 
and later other workers as A.Fisher and V.B.Margquis’ determined the 
potential or metals that have properties that render them difficult 
to handle in water solutions. This method has been applied to tung- 
sten,uranium,and arsenic in particular and while the results that 
have been obtained may not be the best that mignt be desired, they 
are the best that have been brought forward so far. 

These tnen,were the possibilities in tne way of methods that 
had been tried that were open to the author. A study of the chemical 


behavior of chromium had to preceed any attempt at'a determination 


of the potential as well as a general survey of the literature of 


the previous work upon the element in this connection. . 

It has been found tnat chromium can dissolve electrolytically 
in any one of three states, Cro ,Cr Q, Cr O,according to the nature 
of the electrolfye and its temperature. The inactive state of the 
metal corresponds,perhaps, with the passive state of iron. The attai 
ment of this passive state is not to be attributed to the presence 
on the metal of a thin layer of oxide,the hypothesis which is plaus 
ible for iron. It appears rather to imply that the surface,at least, 
of tne metal differs in moleculer arrangement according to the con- 


ditions and surroundings to which it is exposed: in short tnat 


(8) 
chromium appears to be an allotropic element. Chromium exists in 
three different states which are known as active, passive and inter- 
mediate, In each of these it shows a ditferent chemical affinity 


and has a different chemical behavior, 


The selection of proper working conditions necessitated then, 


as mentioned above, an historical treatment of the element, 


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(9) 
HISTORICAL 


a . 
Hittorf found in investigating the electrochemical behavior of 


chromium that in dilute helogen acids or sulfuric acid, chromium 


dissolves on tne application of heat with the formation of cnromous 


compounds, Nitric, chloric, chromic, phosphoric, eitric, tartaric, 


formic, and acetic acids, caustic potash and soda have, however, no 


In its electrical properties it is 


action either hot or cold. 


found to be electronegative, not only to zinc, but also to cadmium, 


iron, nickel, copper, mercury, and silver, and is inactive in 


solutions of these metals except in tnat it reduces mercuric or 


cupric to mercurous or cuprous salts, When employed as anode in 


solutions in whicn it is indifferent, it was found to become covered 


with a yellow film of chromic acid, and tne loss of weignt of the 


anode was found to correspond with the production of sexavalent 


ehromium ions; this occurs even in solutions of hydrogen chloride 


in which chromium ordinarily dissolves with the formation of chromous 


salts, This he supposes to be due either to the decomposition of the 


water by the anion and subsequent formation of chromic acid from the 


lib erated oxygen, or to the formation of a compound of sexavalent 


chromium with the anion and the decomposition of this compound by 


water; no such compound however is known to exist. In solutions 


of potassium thiocyanate or of an iodide, the chromium anode exper- 


liences no loss. A chain .of tne type Cr/KCl/Nano /AgNO /Ag gave no 


electromotive force at 59, and the same result was obtained when a 


dilute acid was substituted for the KCl solution. When the acia 


employed, however, was sufficiently strong to cause the chromium to 


dissolve and tne evolution ot hydrogen, an electro motive force of 


1.056 volts was obtained, When the metal was used for the electrol- 


scien Pa 


OLS Saat 


(10 ) 
ysis of melted KCl or ZnCl , totally different results were obtained, 
and the loss of weight of the anode proved bivalent chromium ions 
were produced, and similar results were found in solutions at 100° of 
the halogen acids,of potassium chloridg,ammonium chloride,and in sol- 
utions of zine and magnesium chlorides at higher temperatures. 

In considering the inactivity of chromium he does notthink the 
theory of the oxide film as applied to iron applies to the case of 
ehromium. In a second paper’ he Suggests this idea. Owing to the 
existence of the different classes of chromium compounds the metal 
may exhibit any potential between the two extreme values;the highest 

alue corresponds with the active state in which chromium gives rise 
to sd ewer compounds. This active state is assumed when the metal 
is placed in sulfuric,oxalic,hydrofluosilicic, or the halogen acids 
at temperatures which are lower the more concentrated the acid; it is 
also active in fused halogen salts or their boiling solutions. 

The metal becomes inactive when immersed in solutions Genesee» | 
free chlorine or bromine,or in strong oxidizing agents. The limiting 

alues, however, last but a short time after removal from the liqnids, 
the metal speedily assuming an intermediate state; the most stable 
active state is that produced vy fused halogen salt. The 
activity is also lowered when the metal serves as an anode for a 
current produced either externally or internally, and with sufficient 
ly strong currents chromic acid is produced, either by decomposition i 
of water by anion and subsequent union of the metal and oxygen, or 
by the intermediate formation of a compound ot the anion and chromium 
When the metal is used as a cathode it rapidly assumes the active 
state. 


9 
A.Fisher in undertaking the study of the electrochemistry of 


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193 Shouse As 66 SeTI5e Cav os ed? wedlw berewoL 


symouslh yo “ed¢le pbeoubosg et biew otmonds a 
Boe legen eft Yo nota i MEupeedue Sas nos 
fico Soe nofae ad? to batroqmn a Te ddtousads $ F 


> @ 


198 Cy BSUS em qibigier #2 obcases # 4e houy sk TAdg 


(7) 

the chromium group or sixth group metals says that the problem is 
complex b~ecause of three main reasons. These are: (P) The “ditty 
culty in preparing aqueous solutions of definite metal content; 
(2) The tendency to show passivity; (3) The preparation of absolute- 
ly pure metal of uniform size and shape. Yisher's paper treats 
only tungsten and uranium, He used the method used by Roderburg 
making use the relation 

E. PF, (water) — E.P. (alecchol) = Constant. 
He found that solutions of WClgin ethyl alcohol showed the least 
passivity and alcoholysis was minimized by suitable dilution and 
maintainance at low temperatures (2°). The element Ww / WCl,in 
ethyl alcohol and LiCl was prepared and compared with Cu / CuCl, 
in ethyl alcohol and LiCl which in turm was compared with Hg / HgClp 
in ethyl alcohol and LiCl, referred to the standard calomel electrod¢ 
on the absolute seale (-0.56 volt). CuCle was selected because of 
its similarity to tungsten inregard to solubility and its nearness 
in the voltage series. By determining the conductivity in the 
alcohol solutions so as to arrive at the value of <,the electrode 
potential in alcohol gave the value for tungsten as 0.680. For 


copper, E P =@.071 volt. From these values 


(water) * etectiol | 
Fisher concludes that for tungsten, E P = 0.680— 0.071 or 


(water ) 
0.509 volts referred to Hp - 0 
In undertaking the study of the electrode potential of chromium 
the literature on the electrolytic preparation of the pure metal 


was also reviewed. 


By using a solution of 100 parts water to 25 parts crcl. and 


working at a temperature of 88° and a current density of 4.3 -5.3 


amperes per square decimeter a good deposit of the metal was 


‘> i a 
ay 


1 ihe 
tf 
A Wind if GD 
ic ace ‘i 


(12) 


obtained by Cowper and Coles? It was found necessary,however, to 


to add an excess of hydrochloric acid to dissolve the precipitated 
oxide and form a clear solution. Solutions of the chromates require 
a high current density which is different for every concentration of 
the chromate and may in no case vary through very wide limits. 

The electrolysis of chromous chlorid shows a side reaction due to 
chlorine. Chromous sulfate is claimed to give very good deposi$s 
of the metal. The sulfate can be prepared by treating potassium 
dichromate with hydrochloric acid; the chromic chlorid obtained 
being rapidly reduced to chromoug chlorideby nascent hydrogen. 

This when treated with sodium acetate gave chromous acetate 

Cr(Ac O)o> which when added in~ excess to dilute sulfuric acid 
constituted the electrolye. Air must be excluded from contact with 


the reacting substances however. 


e1SVOWO!, piSs6aOeA Entiot dex oT *otad baa “iegwod é 


fel igive iq ect avloce zh ‘or bites Site ldoorlyd to seesxs 


VeVESTRSOUOG YIHVS. WOT THeTS TE @2 gy tlw ee Peneb se 


VGdleoto biicia® oimotiie eis aaa tro saris 


“hiss olteFlue eoebib. of es@px<d cz bebba’ gee. 


caiioo suit hetwinrs e¢ seun tla, eyfenspeta 


~ 
~ 


Se «OE ay el snatinenqeneneneneennsenremenmmmemeenas 7 
‘ 
4 


(Si¥ 


Ses ero1ky ess Yo coo ftutog Hoivulce ta8f> a ore 


-.0 Gnls ohéw Q3ev daveisy “anv oes on. af Yeu baw ey 


= 


Ni i o8ss Sh hens ovods Hiro Lis Puccio gS: te R 


hoon y1e* ovina ae bew te Ls et stat low 


« = - 2 - f 


(rack? \¢ Reegel ofleee sbattune eer 


ej a ae 


Sere L de heen wd abis 


oa sta) awongtsc os 


Stagvese epomoide evan etatese bwiboe .di dw 


LONG 


-Tavewvod stoners 


(13) 
EXPERIMENTAL 

In the laboratory of electrochemistry at Illinois zine that 
has recently been amalgamated has been used in experiments on the 
electrode potentials of metals. The results obtained have varied 
so much that there was the desire to improve upon the method. 
Finely divided zine as suggested by Horsch “was then tried. The 
metal was produced by rapid electrolysis of the zine chloride solu- 
tion in which a rod of pure zine served as an anode and a platinum 
wire in the form of a spiral as the cathode. The current density 
was not kept absolutely constant,that is,it varied somewhat with 
successive preparations. According to Horsch this has no apparent 
effect except in the ease of washing the metal. The spongy metal 
so obtained was kept covered with a solution of zine chloride at 
the concentration in which the E.M.F. was to be measured. Upon 
standing for some time a small amount of tne zinc reacted with the 
water to form the hydroxide. This hydrolysis did not effect the 
utility of the zine for after washing with dilute hydrochloric acid 
and then with water, the metal could be used as before. 

The zine chloride solution was prepared by dissolving C.P.zine 
chloride in pure water. Hydrolysis occured to some extent and the 
hydroxide was removed by filtration. The normality of the solution 
was determined by Volhard's method for chiorides. 

The AS-AgCl electrode as prepared by first plating a platinum 
gauze with silver from a cyanide of silver solution and then the 
silver chloride formed as a black deposit by using the silver as an 
anode in a dilute solution of sodium chloride to which a small amount 


of HCl had been added. This electrodé was standardized against 


hydrogen in 0.01 normal HCl and found to have a voltage of 0.4653 


mt NAT, lis Ris Tah i i Ng a Daa a ase Se Sh Racca aac 


(@x) 
~ tadt one: efontl{[{t ts Vitae Ens doors sete to Vidvaroseal . emg 
ett mo €¢dealwegxe nt beer - need sat basames lage cesd bi ks 
‘pe Puy’ eve beaters ai lLues’t ect? ~6lasem To ‘OLalt notg s 
-lorcem ot gege Sorqel of sexlest: sd? ees cred 
bie corty sey" ‘doerven yd bedecanwe 8a ons 
wide eb. ) Canes. SOE Ty Sheytow cele: Bigat: vd 
bres Gh Ona pe eo bevaes eres o1ng “ened 
' (ifenes Fneites Silt  Sboreng ot ae ferige hss 
q fo te tadwougee 76 LTH ney ét teas ties cos ~lesu: 
2"Uaggs on set efit. dsenew (ae untirrenok Bao 


4 uy ie tore emit gat euw to sae eds 
i ofse opts Yo tekoe Lom a ~ at br besevos sqek & 


ei pened of ee: maw 7 «lf, Tt G07. Sotitw oe ol 

. S60-Stiw betdoeer on ts ond 20 Tums Elena se Omty m3 
| ete fon Sib eteylorpys 0 her . oh Enea = 8 
hse of toLde foes siullh Mtiw gaidecw testes: vot. onde 9 
.t0ted es feee 90. binoo Lota oat, 

fit.9.0 poivicesth qd beieqgetg jem notheloe | beta 

bis IiNoixe Gites of Beymno efeyiouier tedine 4 

v3 it to voiteontens aa? Dole ans Let oe . es 
.sebivotdy et Benen 6 1ied lov mee 

alg #-Sabteiaower tt wd hoLeype Ta: aie eborreele & | 

wah Baa Tulos “ev ite To OA LARE ee seo 19vh : 

ga "evlis ad3 se tttes “e oP Leangeb doa ly: & BS, beano. 4 
peeone Lous « iocdw oF ebivalad wutbow to notvisfow sat 
fefiicse fee tinwbpasa sil Sheree BLA ysbobha 
Eae.0 to sya lov 2 ova og bee? baa £98 amie. Jot 


(14 ) 
volts,this reading being the average value obtained from a number. 


The complete cell. was then set up as shown on the following 


page. It consisted of the chain Zn/ZnC1,/ AcOlfvAg . A table of 


E.M.F. readings is given below and represents six eells as numbered. 


(1) (2) (3) Cell Number (4) (5) (6) 
Poses. 1.1580. 1.1600 ieLogo 11649, 151580 
Pmoets tooo, lkl6LO bel604) Veer 2 ie7e 
iolece £.2S50'° L610 COY™ Tilese Wi .Lsa2 

These results show, in a measure, the very thing that it was 
desired to correct in the method as used in the class work. While 
only the values for 0.01 molal solution are given, the author is 


of the opinion that the data are enough to prove the uselessness of 


trying to use this as a class experiment in this form. 


Experimental (Con't) 
Work Done Upon Chromium in the Laboratory. | 
12 
It was hoped that in a manner similar to that used by A.Fisher 


a salt of chromium might be prepared which would be soluble in 


ethyl alcohol and then by means of the constant E.M.F. existing 
between the value for alcohol and water the electrode potential of 
the element arrived at. With this idea in mind the author prepared 
some anhydrous chromic chloride. 

Metallic chromium was placed in a pyrex tube about three feet 
long and by passing dry chlorine over the red hot metal for some 
time a product was obtained. The chlorine was geherated by drop- 
ping concentrated hydrochloric acid upon crystals of potassium per- 
manganate. The gas was washed by passing through two bottles of 


water and then a third bottle of sulfuric acid which served to 


remove the water. A U-tube containing calcium chloride was also 
placed in the chain for the same purpose. The metal glowed strong- 


Sef up as used 
im aererrmitrg electrode potentials. 


STANDARD 
CELL 
| BATTERY 


are 


TAPPING 
KEY 


eSC0S9 0 B00 


POTENTIOMETER 


Calomel Sketches Half-cell 
Half- cell — of — for metals 


+ S¥Stt ‘ \o4 .< ’ <q - to. ——- 


yes att er et 


Seu) 29 gy tae 
alorinstog SHON IaIS QOMTIeH—D, | 


it90-410Y4 aor} sane 


~ 


? 


(15) 
ly after the reaction had begun showing that it was endothermic. 

After cooling,the contents of the tube were removed and found 
to consist of a mixture of products. A fine apple blossom colored 
powder that had sublimed away from the rest of the mass and a large 
bulk of purple crystals containing some of the metal intermixed 
with them,were obtained. From discriptions found in the literature 
it was decided that both the crystalline mass and the fine powder 
were chromic chloride. 

It was found that the chromic chloride was insoluble in alcoho. 
and hence this line of endeavor reached an end. A more soluble 
salt was sought after. There appeared to be no chromic salt that 
was easily prepared with the means at hand suitable for the purpose. 
In the literature a method for preparing chromous chloride from 
the chromic salt is discribed. 


The purple crystals discribed above were placed in a pyrex 


tube and it was hoped that by passing hydrogen over them in a heated 


condition they would be reduced to the easily soluble chromous 
chloride. The hydrogen was obtained from a cylinder of electrolyt- 
ic hydrogen. Even strongly heating the chromic chloride with tne 
gas did not give a quantity of the chromous salt worth separating. 
There was some to be sure, but it was intermingled so thoroughly 
with the dark mass of residue that it was impossible to obtain even 
a sample of the salt without hoping to be able to work with it. 

A repetition of the operation using a fused quartz tube in order 
that higher temperatures might be reached, yielded the same kin a 

of resulte. 


A chromium electrode was prepared by electrolyzing a solution 


of chromic chloride with a current density of about four amperes 


‘ 
> witane 

Io LS VES 4 
es 

‘ne? 


a 


(16 ) 
per square centimeter. The chromium was plated upon a piece of 
platinum foil but was not very adherent. That is, due to some 
defect in the current density or the temperature the deposit was 
crystalline instead of the smooth coating that it had been hoped 
would be obtained. This defect,it is thought, could have been cor- 
rected had there been any real need for the metal. Since a suit- 


able electrolyte in which to make the determination had not been 


found the metal could not have been used as an electrode even, if 


it had have been obtained in the form of an even smooth coating. 


_-~ «6 “yu 
hs on wean ; a 
i oti ' i. * ; i 
<5 et, a ‘- 
— _ ” i a aa pre = tee cette ee — J 


tec eoelg & noqy bevale Saw i isons sft " Neltnoosdis 
emoe oF sub ,et tadT  .castedha yrey for ean dtd Tip 
uw tieoqeh ent s«wisvegmed’ edd 40 ‘it derteh Saetuis ost 
‘eqodt med bad ¢i Jad} gaitaon: droome ens to baséual 
70m «need evad Sivoo tidguodt e2 Jf Josteh wick begt 
ina @ soatt - Lasem on? 19Y bsen Lees ye Mt 
66S con fed dtendiaieden eft eaten Oo. ao it 
abertesls am se heey ceed evad son 


~BCiseos Sl00ms E9978. he 14 cot eit af 
: res 


id 
p 
© 


cA 


* 

i 

$y ies 
Sen = 


1 
: Owe aes 
” P ry 


& ; . 
Me 
f ¥ 
baat Tad 
. ‘Vi 
> Hy 
[ ‘ 
+ *, v 
gia 
® sa 
‘ Ne - Mi 
1, Ser ae nt. é; 
; ae a oe: ‘ 
b io? - a a oie be a a eke acy \ — 
* ; my - 7 Bee ss 
: i y ike; as ete ue nae , 
- ee eee - va a ‘© ee 


ee ees Me His oh ae ad meV Ke Ps tas 


A SUMMARY OF THE 
EXPERIMENTAL RESULTS OBTAINED. 


Finely divided zine was prepared by the electrolytic deposition 
or the metal upon a platinum wire from a solution of the chloride. 

Using the chain Zn/ZnC1,/AgC1/Ag results were obtained which 
did not seem to warrant the use of the experiment as a class exper- 
iment. 

It was found that the method of Roderburg wherein tne difference 
of E.P. between absolute alcohol and water is made use of,was not 
applicable in the case of chromic chloride since it is not soluble 


in alcohol. 


An attempt to make anhydrous chromous chloride was a failure. 


ee ee OO re 


SS me RT 


-Shftoins efy *o nclizvios s BOS vl minh tate @ “a 


Cewisstih any cheredtu: guuttteboe *o totes ‘$n sake. 


oS cereale 5p mee? eam 


———— -—~ Ci be a a A ap pnp ado al 


(Fz) 


EET To Yaamue * | 
ORMIATHO GtIvess TAT WArLANeE 


A : 
~ 
= t. 


teiveqes cisyfoutosts ont ge cereae. ve mts 4 


tw bentesdo s16w, ed Ses SA \ COA Ng ADAG \a8 er ay) 


~~; 


-s9Gxe seals « sa Tremizeqxa ai? Do eae, 998 inert 


——— 


he 


Pon eaw,to séy Gham oe! Yedsw dae fodoo Le erulosea, ae 
éidufoe jon el cs eonis pikrelio plnovie to —_ x0 08 


! fe +e 


-Stelict @ ssw ebivefdo evomonds exothydas odam off 


on - — en aeiiee ahesehiaemeel aa 


BIBLIOGRAPHY 


----- G.Horsch, Journal Am.Chem. Society-41,1787 (Nov.1919) 


----- Kahlenburg, Trans .Am,Elect .Chem.Soc .36,277(1919) 


j----- Noyes « Keyes. J.A.C.S. Vol .35,340. 
(essse Pyys.Review. 6, 211-282 (1915) 

D----- A.Fisher.,Z.Anorg .Chem.35,170-200. 
6----- W.B.Marquis.,J.A.C.8.42,1569.(1920) 
?----- Hittorf,Zeit.Phys .Chem.25,729-749.(1893 ) 
eSSs= Pi ttert,, * ud “sl 30 ,481-507 .(18y9y) 
SSSSSe Zeit.Anorg.Chem. 41,170-20xu. 


10----Cowper « Coles. Chem.News.o1,16, (1900) 


it----@.-Heorscn,J .4.€.9. 41,2787 .(1919) 


12----A.Fisher,Z.Anorg .Chem.35,17/0-208. 


- 


RRNA ee lt ae 
- a + ss 


ACKNOWLEDGMENT 


The author wishes to take this opportunity to express his 
appreciation of the many helpful suggestions and criticisms received 
from Dr. Gerhard Dietrichson under whose direction this work was 


earried out. 


ETA 
EEE 
EEEE="- 


———— 


